SnailsAttack
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Acid exchange reactions
| Code: | Prof. -----------,
Do you know anything about acid exchange reactions? Like, the ones where there's a double displacement between an acid and a salt? I'll cite some examples I've found on the internet:
= Hydrochloric acid from sulphuric acid
H₂SO₄ + NaCl -> HCl + NaHSO₄ //supposedly occurs at room temperature in aqueous solution
2NaHSO₄ + 2NaCl -> 2HCl + Na₂SO₄ //supposedly requires mild heating in aqueous solution
= Nitric acid from sulphuric acid
H₂SO₄ + NaNO₃ -> HNO₃ + NaHSO₄ //supposedly occurs at room temperature in aqueous solution
2NaHSO₄ + 2NaNO₃ -> 2HNO₃ + Na₂SO₄ //probably chemically feasible through mild heating in aqueous solution
= Nitric acid from hydrochloric acid
HCl + NaNO₃ -> HNO₃ + NaCl //supposedly occurs at room temperature in aqueous solution
= Acetic acid from sulphuric acid
H₂SO₄ + NaCH₃COO -> CH₃COOH + NaHSO₄ //supposedly occurs at room temperature in aqueous solution
H₂SO₄ > HCl > HNO₃ > CH₃COOH (reactivity series?)
Sulphuric acid seems to be the king of acids when it comes to displacing weaker acids from their salts. Do you know why that is? Do acids have a reactivity series the same way metals do? If that's the case I can't seem to find much about it online.
In short, I want to know how to predict the outcomes of the following acid-salt interactions:
H₃PO₄ + 3NaCH₃COO <-?-> 3CH₃COOH + Na₃PO₄ //acetic acid from phosphoric acid (or vise versa?)
CH₃COOH + NaNO₃ <-?-> HNO₃ + NaCH₃COO //nitric acid from acetic acid (or vise versa?)
6CH₃COOH + 3Na₂B₄O₇ <-?-> 2H₃BO₃ + 6NaCH₃COO //boric acid from acetic acid (or vise versa?)
Any insight would be appreciated.
Thanks,
------ ---- |
This is an email I wrote to my chemistry professor about a week ago. Their response was... not enlightening.
These "acid exchange"(?) reactions I'm talking about are not the same as typical salt metathesis reactions which proceed by evolution of an insoluble
precipitate. Rather they seem to be controlled by some sort of acid strength hierarchy.
Tell me what you guys think.
[Edited on 2/10/2022 by SnailsAttack]
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SWIM
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Many of these reactions proceed because the product is a gas which escapes, either pure or along with water, driving the equilibrium forward.
Hydrochloric acid is, I think, stronger than sulfuric.
If the hydrochloric isn't being boiled off then you've got sulfate and chloride ions mixed up with sodium and hydronium ions as both of these acids
are fully ionized in dilute solutions.
Kind of hard to say what's what in such a situation.
Edit: edit deleted. I see what that was about.
Amazing what you miss when you don't participate in social media.
This site is about as social as I get on the internet.
Another edit: rainwater explains it better, even with part of his post missing.See below.
[Edited on 10-2-2022 by SWIM]
[Edited on 10-2-2022 by SWIM]
[Edited on 11-2-2022 by SWIM]
[Edited on 11-2-2022 by SWIM]
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Rainwater
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Ya'll check my work here. This is an easy one. Im going to try without looking it up
H₂SO₄ + NaCl -> HCl + NaHSO₄ //supposedly occurs at room temperature in aqueous solution
So first, let's add the ac8d to the water
H2O + H2SO4 = H3O+ + HSO4-
Or simply
H+(aq) + HSO4-(aq)
Then we add the salt.
Which in aqueous solution becomes Na+ and Cl-
H+(aq) + HSO4- + Na+ + Cl-
Mr attacks. Please read this while i ramble about stuff i just learned the other day and some stuff i made up to help myself understand
table of acid bace strengths
So we have dissociated a bunch of stuff and have catons and anonds floting around in solution. Different things have different attractions to protons.
H+ is just a proton. So unless something like temperature, pressure or the loss of items from the solution happens. Then the stronger
attraction wins the proton.
In this case Cl- has a ka of 1300000 and HSO4- has a ka of 1000
Forgot t9 added. Ka is how easily something loses its protons
Here is where i go crazy and prove that i know nothing about chemistry
Edit.
Lost half my post
So, as of right now, no chemical reaction has taken place.
Just all this stuff floating around. As the conditions change, you will have a reaction. Say the solution starts evaporating. Once the consentrations
get large enough, ions will fall/float out of solution . Sodium salt and hydrochloric gas
Lets simplify this.
Dont use aqueous solution
Pure Sulfuric acid and salt.
Now, as the ions dissociate, there will be a reaction.
Hcl isn't very soluble in sulfuric acid.
You will see bubbles excaping the solution. Hcl is a gas at Stp. So, it will leave the solution driving the reaction forward.
[Edited on 11-2-2022 by Rainwater]
[Edited on 11-2-2022 by Rainwater]
[Edited on 11-2-2022 by Rainwater]
"You can't do that" - challenge accepted
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Texium
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Mmm I’m passionate about acid-base chemistry.
All Brønsted acids (which can actually be anything containing hydrogen, just depends on the context) have an equilibrium constant, known as Ka. A
more practical way to look at these is as the pKa, which is the pH at which the concentration of the acid (HA) and the dissociated form (H+ A-) are
equal. Effectively, at a pH above the pKa of an acid, it will be deprotonated, and at a pH lower than the pKa, it will be protonated. In aqueous
solutions, the pH of the solution is limited by the pKa of water, which is 14, and the pKa of the hydronium ion, which is -1.74. Thus the typical pH
scale seen for aqueous solutions of about -1 to 14. If you add an acid with a lower pKa than hydronium to water, it will fully dissociate and
protonate the water to create the less acidic hydronium ion. That’s the definition of a strong acid (for example, hydrochloric acid has a pKa of
-6.3). So consider you have a 1 M solution of HCl (which, pedantically would really be a solution of hydronium and chloride ions in water, with a pH
of 0, well above HCl’s pKa of -6.3). Then you add sodium acetate to that solution . Acetic acid’s pKa is 4.75. The pH of the solution is 0- much
lower than that. So the acetate gets protonated to acetic acid. As long as there is still an excess of HCl to keep the pH well below 4.75, most of the
acetate will be converted to acetic acid which will not dissociate at that pH.
So in summary, yes, there is a hierarchy of acid strength, and you can easily find it by looking up a pKa table. There’s a lot out there, some of
which are simple, and others of which cover everything from superacids well into the double-digit negatives to methane at +50. These values are
extremely useful both for inorganic and organic chemistry!
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DraconicAcid
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To add to what Texium says, you can consider an acid-base reaction to be a competition between two bases for H+.
HX + Y- <===> X- + HY
If Y- is a stronger base than X-, then this equilibrium will lie to the right (K > 1). If X- is more basic than Y-, then equilibrium will lie to
the left (K < 1). The strength of an acid and its conjugate base are inversely proportional- if Y- is a stronger base than X-, then HY is a weaker
acid then HX.
There are a number of factors which determine how basic a species is- its charge, the electronegativity and size of the atom with the lone pair, the
presence of electron withdrawing groups, and the delocalization of electron density through resonance.
Hydroxide ion is a strong base- it's very good at taking H+ away from other species. It's conjugate base is water, which is as weak as an acid can be
and still be considered an acid in aqueous solution. Any base that is stronger than hydroxide (such as oxide, amide ion, or hydride ion) cannot exist
in water (it will immediately react to give hydroxide); the conjugate acid of any base stronger than hydroxide will be a feeble acid (hydroxide,
ammonia and hydrogen gas), and have no acidic properties in aqueous solution.
Bases that are weaker than hydroxide (such as acetate, cyanide, or fluoride) will have conjugate acids that are more acidic than water (acetic acid,
hydrocyanic acid, or hydrofluoric acid).
Strong acids (such as HCl, HNO3, HClO4) will also not exist in aqueous solution- they will ionize completely to give hydronium ion and their conjugate
bases. These conjugate bases (chloride, nitrate, perchlorate) are feeble bases, and will have no basic properties in aqueous solution.
The pKa of hydronium ion is often quoted as -1.74, but that always struck me as wrong. The Ka for hydronium is the equilibrium constant for the
reaction H3O+ + H2O == H2O + H3O+, which is obviously a null reaction, with a K = 1. Calculating pKa as -1.74 requires that we ignore water on one
side (because it's the solvent), but NOT ignoring water on the other side (because it's the conjugate base of hydronium)- an artificial distinction at
best.
In concentrated acid solutions, pKa values fall apart, because they only really apply in dilute aqueous solutions. In a dilute solution, you can't
react sulphuric acid with sodium chloride to get any hydrochloric acid, because all you have is hydronium ion. In concentrated sulphuric acid, you
can protonate the chloride ion and distill off hydrochloric acid (because HCl is much more volatile than sulphuric acid).
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Rainwater
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Poor mans aqua regia as i have heard it called.
Hcl come in aqueous solution. I get it in 33.7%(w/w) so its
33.7% hcl gas dissolved in water.
H+ + Cl- + Na+ + HNO3-
It can be seen as HNO3 and NaCl
But how do you purify any one compound.
If we could make a list of all possible compounds that have these ingredients and list their melting points, boiling point, and solubility that would
help us.
We see that 37% hcl(aq) boils at 48c. 10% hcl(aq) at 103c
Nitric acid has a 120c azeotrope so distillation would not produce a good sample
Even if we reacted HCl gas directly with the salt without water, this would be an issue because sodium cloride is soluble in nitric acid.
Lets look at fractional freezeing.
Hcl @ 33% melts at -36c
HNO3 melts @ -42c
So freezing will not work ether.
Lets get mad and toss a copper penny into the solution.
Now, you have a reaction going. As the nitrate ions react with the copper metal to produce a gas and a percipitate, just remember that nitric acid
also reacts with trousers. And orange gas is usually toxic
"You can't do that" - challenge accepted
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DraconicAcid
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Trying to make nitric acid from hydrochloric acid (and separating it from the mixture) seems like a fool's errand to me.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Rainwater
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Correct. If you need pure hno3(aq) you will not get it using hcl.(atleast no way i know of)
H₂SO₄ + KNO₃ -> HNO₃ + KHSO₄ is my go to. Distilled under -15inhg at 80c.
So KHSO₄ doesn't boil until +300c(decomposes), and the nitric acid(+95%) boils at 80c under reduced pressure. So, the components are very easy to
separate.
Edit:
Just checked my log book. Not sure about the vacuum i quoted above. Looks like when i used tap water in my aspirator it boils between 78-80c. Ice
water 63-67c.
I know when using ice water i can boil pure h2o at 35c. Not sure how to figure the boiling points under different pressures yet.
[Edited on 11-2-2022 by Rainwater]
"You can't do that" - challenge accepted
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Texium
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Look up the Clausius-Clapeyron
equation.
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ave369
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Let me add my two cents.
The acid strength hierarchy is important, but there is another factor, called the Le Chatelier principle. It says that the equilibrium of a chemical
reaction is shifted if you change concentration of one of the reagents, temperature or pressure. Its most important consequence is that if we
constantly remove one of the reagents from one of the sides of the reaction, the equilibrium will shift to that side.
For example, if one of the acids is volatile and the other is not, we can distill the mixed solution and get the volatile acid to evaporate. The
nonvolatile acid will react further, until the reaction is driven to completion. Thus you can use a weaker but nonvolatile acid (e.g. phosphoric) to
distill a stronger but more volatile acid (e.g. hydrobromic).
The same applies if one of the acids is unstable and decomposes (e.g. carbonic or ferric acid). As the unstable acid decomposes, the equilibrium is
shifted towards making more of it.
And now the answer to why sulfuric acid is the mother of all acids: it is really strong, really stable and really nonvolatile. That's why it displaces
nearly any other acid from its salts.
[Edited on 13-2-2022 by ave369]
[Edited on 13-2-2022 by ave369]
Smells like ammonia....
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RustyShackleford
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Quote: Originally posted by ave369  | Let me add my two cents.
The acid strength hierarchy is important, but there is another factor, called the Le Chatelier principle. |
more important than pka, oxalic acid is able to form nitric acid in 95+% conversion (determined by yield of nitric, likely even higher) from calcium
nitrate. Since calcium oxalate is incredibly insoluble, and theres no mechanism by which it would redissolve, the oxalic is able to form the many many
times stronger nitric acid in near quantitative yield.
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SnailsAttack
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thanks a ton you guys. i always figured that all acids fully dissociate in solution the same way as s̶a̶l̶t̶s salts of strong acids. the concept
of pH and acid equilibriums were kind of lost on me in highschool as "just a bunch of numbers" (and they might still be) but i'm starting to get it
now.
here's some notes i took from what you all said
| Code: | - the reason that hydrochloric acid, nitric acid, and acetic acid can be produced using sulphuric acid is because those acids all have a lower boiling point (there's azeotropes involved for many of them, but that's the main idea at least).
- acid 'strengths' are given by this table: https://depts.washington.edu/eooptic/links/acidstrength.html
- Ka is the acid dissociation constant. it's essentially a measurement of how easily an acid is deprotonated. pKa is the logarithmic form (pKₐ = -log₁₀Kₐ), which can also be defined as the pH at which the concentration of the acid (HA) and the dissociated form (H+ A-) are equal.
- at a pH above the pKa of an acid, it will be deprotonated, and at a pH lower than the pKa, it will be protonated.
- in aqueous solutions, the pH of the solution is limited by the pKa of water, which is 14, and the pKa of the hydronium ion, which is -1.74. Thus the typical pH scale seen for aqueous solutions of about -1 to 14. (i've seen 0 to 1 but supposedly it's technically -1.74 to, like, 15)
- if you add an acid with a lower pKa than hydronium to water, it will fully dissociate and protonate the water to create the less acidic hydronium ion. that’s the definition of a strong acid.
- in concentrated acid solutions, pKa values fall apart, because they only really apply in dilute aqueous solutions. (pH is sort of a manifestation of the interaction of acids with water)
- nitric acid can't be distilled using hydrochloric acid because their boiling and freezing points make separation impractical. this is why sulphuric acid is used.
- the acid strength hierarchy is important, but there is another factor, called the Le Chatelier principle. It says that the equilibrium of a chemical reaction is shifted if you change concentration of one of the reagents, temperature or pressure. Its most important consequence is that if we constantly remove one of the reagents from one of the sides of the reaction (usually this will be by distillation), the equilibrium will shift to that side.
- oxalic acid is able to form nitric acid in 95+% conversion (determined by yield of nitric, likely even higher) from calcium nitrate. since calcium oxalate is incredibly insoluble, and theres no mechanism by which it would redissolve, the oxalic is able to form the many many times stronger nitric acid in near quantitative yield. |
so far i've only worked with vinegar. this'll be useful as i work my way towards some of the more exotic acids. might try for citric next if i can
come up with a good recrystallization process
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Rainwater
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Im stilling your notes. Great summary
Extracting it from fruit is a difficult and cheap exercise in purification. Great starting point. Best I was able to get was %14 theoretically yield
from concentrated lemon juice.
Red nile on youtube has a video demonstrating the process. Good tutorial. Much cheaper / easier / purer to buy it.
https://youtu.be/FMtayizdFiw
"You can't do that" - challenge accepted
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clearly_not_atara
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The reason that you're not getting a simple answer is because there isn't a simple answer. There are several processes at work.
- The proton dissociation equilibrium determines the number of undissociated molecules of an acid in solution as a proportion of the total number of
molecules.
- Raoult's Law of Volatility determines the probability for any component of a mixture to evaporate
- The product of concentrations of components of a salt determines the probability that the salt will precipitate
- Changes in the concentration of components of a solution causes other components to react to maintain the equilibrium according to Le Chatelier's principle
- Concentrated acids, i.e. not aqueous, can act as though they are much stronger than they would be in aqueous solution. NaHSO4 +
NaCl can produce HCl not because of any of the other principles, but because this reaction occurs where NaHSO4 makes up a large fraction of the
"solvent", effectively molten sodium bisulfate monohydrate. H3PO4 + KI >> HI is also like this. In general, if the mass fraction of an acid (or
alkali) is greater than 10%, the usual quantitative rules start to bend quite a bit.
One interesting consequence is that oxalic acid can produce dilute solutions of sulfuric acid by reaction with magnesium sulfate, because magnesium
oxalate is highly insoluble. But concentrated sulfuric acid will displace magnesium oxalate.
[Edited on 16-2-2022 by clearly_not_atara]
[Edited on 04-20-1969 by clearly_not_atara]
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